Le Chatelier's principle
Henry Le Chatelier was a French chemist who studied reversible reactions and equilibrium and he made some
Equilibrium is the point in a reversible reaction where the rate of the forward and reverse reactions are the same.
We can show this as:
During a reversible reaction the reactants turn into products and the products back into reactants again. The
important point to remember is that these two reactions occur at the same time. The forward reaction will proceed at
a given rate, labelled Rf and the back reaction which turns products back into reactants will also proceed at a given rate, labelled
Rb in the example above. The rate of the forward and reverse reactions can be altered if we change the reaction conditions, e.g. for
example if we alter the concentration of one of the reactants or products or change the temperature of the equilibrium mixture, but
recall that at equilibrium the rate of the forward reaction, Rf and
the back reaction, Rb will be the same.
Equilibrium is a very stable low energy point for a reaction, at
equilibrium the reaction is nice
and happy!! It's like the ball at the bottom of the curve. If you push on it so that it
moves away from the bottom of the curve the ball will roll back down again, to the low energy point.
Well chemical reactions are like that- if a reaction is at equilibrium (rate of forward and
reverse reactions are the same) and you come along and heat it up or put it under
reaction will re-adjust itself to get back to a new equilibrium position, that is a low energy state.
At this new equilibrium
position the amounts of reactants and products maybe different but the rate of the forward and reverse
reactions will be the same e.g. Consider the reaction of nitrogen and hydrogen to make ammonia, the Haber Process.
Nitrogen(g) + hydrogen(g) ⇌ ammonia (g)
N2(g) + 3H2(g) ⇌ 2NH3(g)
At equilibrium the mixture contains only a small amount of ammonia and is mostly nitrogen and hydrogen.
We would say the position of equilibrium lies to the
left (reactants). Since ammonia is a useful substance we need to shift the position of
that it moves more to the right (products).
The position of equilibrium for a reaction depends on several variables, these include concentration,
and pressure (if gases are involved in the reaction).
We can adjust the amounts of reactants and products at equilibrium by simply changing the reaction conditions.
The trick is to know exactly what conditions to change and that is exactly what Henri Le Chatelier figured out.
Le Chatelier's Principle is the idea that a system (reacting chemicals) at equilibrium will
oppose any changes applied to it. We will use the Haber Process to try and explain Le Chatelier's principle in more detail.
The equation for the Haber process is shown again below:
Nitrogen(g) + hydrogen(g) ⇌ ammonia (g)
N2(g) + 3H2(g) ⇌ 2NH3(g)
The forward reaction:
N2(g) + 3H2(g) → 2NH3(g)
Is exothermic, that is it releases heat energy to the surroundings. The back
2NH3(g) → N2(g) + 3H2(g)
Is endothermic, that is it removes heat energy from the surroundings.
So if you have a mixture of nitrogen, hydrogen and ammonia in a flask at equilibrium. The position of
equilibrium for this reaction lies very much to the left and unfortunately there is very little
ammonia in the
flask. The challenge for chemists is how to adjust this equilibrium so that the amount of the
product ammonia can be increased.
- What would happen if you warmed the flask? Well according to
Le Chatelier's principle the reaction
should oppose any change you make to it, that is it will try and remove the heat you added. How would the added
heat be removed? An endothermic reaction will
remove heat! The
only way to do this is to force the equilibrium position even more to the left, that is to increase
the back reaction since this is an endothermic reaction:
2NH3(g) → N2(g) + 3H2(g)
- However this is bad news as it will reduce the
amount of ammonia in the flask. So to increase the amount of ammonia instead of
heating the flask, cool the mixture in the flask.
Again according to Le Chatelier's Principle the system will try to generate
heat, to oppose the
cooling, so it will force more nitrogen and hydrogen to react to make ammonia since the forward reaction
is exothermic. However we know that cooling a reaction will slow it down. So
ultimately cooling it down will increase the amount of ammonia in the flask but it might take a very long time for the new equilibrium
position in which there is more ammonia present to become established.
- What would happen if you increased the pressure to the position of equilibrium in the Haber process?
Well first thing is that pressure will only have an effect if there are gases involved in the reaction. So since all the chemical reacting here are gases then
will have an effect.
At room temperature and pressure 1 mole of any gas will occupy 24 litres. So for the reactants the
total number of moles is 4 so this will occupy a volume of 96 litres. For the products we have
2 moles of gas, so this will occupy 48 litres at room temperature and normal pressure
(standard temperature and pressure S.T.P). So obviously we can think of the reactants as the
high pressure side of the equation and the products as the low pressure side. So if you increase
the pressure on this reaction then it will force the equilibrium to the right, that is the
equilibrium mixture will contain more ammonia. If you reduce the pressure then the equilibrium
mixture will contain more reactants and less ammonia.
- Similar arguments are true if you add more reactants or products then the system will no
longer be at equilibrium and it will adjust the concentrations to achieve a
new equilibrium position.
If you add more reactants it will force the equilibrium to the right, that is towards the products,
similarly if you add more products it will force the equilibrium towards the left side, that is the
amount of reactants will increase.
- How about adding a catalyst? Catalysts do not have any effect on the position of equilibrium,
they will not affect the amount of reactants or products in the equilibrium mixture. What a catalyst
will do is increase the rate of both the forward and reverse reaction in an equilibrium mixture.
The catalyst will allow the system to achieve equilibrium faster.
Dinitrogen tetroxide (N2O4) is colourless toxic gas with a very unpleasant smell. Despite the fact that
it is a colourless gas it appears reddy brown since it exists in an
equilibrium mixture with brown nitrogen dioxide gas. This can be shown as:
nitrogen dioxide(g)⇌ dinitrogen tetroxide(g)
2NO2(g) ⇌ N2O4(g)
The forward reaction, the conversion of nitrogen dioxide into dinitrogen tetroxide is
exothermic which means that the reverse or back
reaction is endothermic.
What would happen to the position of equilibrium if we:
- placed a flask of this equilibrium mixture of gases in a beaker of iced water? Well
according to Le Chatelier's principle
the reaction should oppose any changes made to it. So in this case the change is a reduction in temperature
so the position of equilibrium
will shift to the right, that is make make dinitrogen tetroxide since this reaction is
exothermic and will generate heat energy. This would mean the colour of the
equilibrium mixture of gases in the flask should fade, since the
concentration of the brown gas nitrogen dioxide is reducing.
- Using the opposite reasoning if we were to place a flask containing an equilibrium mixture
of these two gases in a
hot oven, what would happen to the position of equilibrium? Well according to
Le Chatelier's principle the system will try
to oppose the temperature rise, so it will push the equilibrium in the direction
of the endothermic reaction, that
is there will be more nitrogen dioxide in the
equilibrium mixture. This means the flask will get darker. This is shown in
the image oposite.
Le Chatelier's Principle and pressure
Using Le Chatelier's principle to explain changes in the position of an
equilibrium due to changes in pressure
only applies to
reactions involving gases. Solids and liquids are not affected by changes in pressure.
As an example consider the reversible reaction that occurs between the colourless gas nitrogen tetraoxide (N2O4)
and the brown gas nitrogen dioxide (NO2).
What about if we had the gases in a syringe and we compressed the syringe? Well this would increase the pressure. So what would happen
this time to the concentration of gases at equilibrium?
You can see from the equation opposite that 1 mole of the colourless N2O4 gas
dissociates to form 2 moles of the brown
NO2 gas. Since there is 1 mole of gas on one side of the
equation and 2 moles on the other we can say that the side with 1 mole of gas
will be the low pressure side and the side with 2 moles of gas will be the high
pressure side. The image
below shows what happens to
an equilibrium mixture of N2O4 and NO2 gases in a syringe as the
pressure is reduced and increased:
- A reaction at equilibrium is one where the rate of the forward and reverse reactions are the same.
- Le Chatelier's principle explains what happens to reactions at equilibrium. According to his principle a reaction at
equilibrium will change or alter the position of equilbrium in order to minimise or cancel out the change applied.